Chapter 1

Prentice Hall © 2003 Chapter 6

Chapter 1Chapter 1Atoms, molecules, & chemical Atoms, molecules, & chemical

bondingbonding

CHEMISTRY

David P. WhiteDavid P. White


The Discovery of Atomic The Discovery of Atomic StructureStructure

• The ancient Greeks were the first to postulate that matter consists of indivisible constituents.

• Later scientists realized that the atom consisted of charged entities.



Cathode Rays and Electrons


Cathode Rays and Electrons• A cathode ray tube (CRT) is a hollow vessel with an electrode at either end.

• A high voltage is applied across the electrodes.

• The voltage causes negative particles to move from the negative electrode to the positive electrode.

• The path of the electrons can be altered by the presence of a magnetic field.

• Consider cathode rays leaving the positive electrode through a small hole.

– If they interact with a magnetic field perpendicular to an applied electric field, the cathode rays can be deflected by different amounts.




Cathode Rays and Electrons– The amount of deflection of the cathode rays depends

on the applied magnetic and electric fields.– In turn, the amount of deflection also depends on the

charge to mass ratio of the electron. • In 1897, Thomson determined the charge to mass ratio of

an electron to be 1.76 108 C/g.


Cathode Rays and ElectronsMilikan’s oil drop experiment



Milikan’s oil-drop experimentConsider the following experiment:• Oil drops are sprayed above a positively charged plate

containing a small hole. • As the oil drops fall through the hole, they are given a

negative charge.• Gravity forces the drops downward. The applied electric

field forces the drops upward.• When a drop is perfectly balanced, the weight of the drop

is equal to the electrostatic force of attraction between the drop and the positive plate.



Milikan’s oil drop experimentUsing this experiment, Millikan determined the charge on

the electron to be 1.60 10-19 C.

With more accurate numbers, we get the mass of the electron to be 9.10939 10-28 g.



RadioactivityConsider the following experiment:• A radioactive substance is placed in a shield containing a

small hole so that a beam of radiation is emitted from the hole.

• The radiation is passed between two electrically charged plates and detected.

• Three spots are noted on the detector:– a spot in the direction of the positive plate, – a spot which is not affected by the electric field,– a spot in the direction of the negative plate.



Radioactivity



Radioactivity• A high deflection towards the positive plate corresponds

to radiation which is negatively charged and of low mass. This is called -radiation (consists of electrons).

• No deflection corresponds to neutral radiation. This is called -radiation.

• Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called -radiation.


The Nuclear Atom• From the separation of

radiation we conclude that the atom consists of neutral, positively, and negatively charged entities.

• Thomson assumed all these charged species were found in a sphere.



The Nuclear Atom• Rutherford carried out the following experiment:• A source of -particles was placed at the mouth of a

circular detector.• The -particles were shot through a piece of gold foil.• Most of the -particles went straight through the foil

without deflection.• Some -particles were deflected at high angles.• If the Thomson model of the atom was correct, then

Rutherford’s result was impossible.



The Nuclear Atom• In order to get the majority of -particles through a piece

of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge the electron.

• To account for the small number of high deflections of the -particles, the center or nucleus of the atom must consist of a dense positive charge.

The Nuclear Atom• Rutherford modified

Thomson’s model as follows:– assume the atom is spherical

but the positive charge must be located at the center, with a diffuse negative charge surrounding it.

The Discovery of The Discovery of Atomic StructureAtomic Structure

• The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons).


The Modern View of The Modern View of Atomic StructureAtomic Structure


The Modern View of The Modern View of Atomic StructureAtomic Structure

• Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus.– There can be a variable number of neutrons for the same

number of protons. Isotopes have the same number of protons but different numbers of neutrons.

• Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons.

The structure of the atomThe structure of the atom

• Matter consists of very small, indivisible particles, which are named atoms

• Atoms are made up of subatomic particles- p, n & e– Electron: negative charge,

– Theories about the energy and the arrangement of electrons in atoms are based on the interaction of matter with electromagnetic radiation.

– dual nature: wave & particle



• All waves have a characteristic length, height & number of waves that pass through a certain point in one second

• wavelength, • amplitude, A.• The frequency, • The frequency of electromagnetic radiation is related to

its wavelength by:

The Wave Nature of LightThe Wave Nature of Light

c


• Electromagnetic radiation moves through a vacuum with a speed of 2.99792458 10-8 m/s.

• Electromagnetic waves have characteristic wavelengths and frequencies.

• Example: visible radiation has wavelengths between 400 nm (violet) and 750 nm (red).


Example Example

• The wavelength of the green light from a traffic signal is centered at 522nm. What is the frequency?


c

1149

8

1075.510522

/1000.3

sm

smcv

• Heated solid emits radiation over a wide range of wavelength

• Dependence of amount of radiation energy emitted on wavelength- partially successful with wave theory


• Planck: energy can only be absorbed or released from atoms in certain small amounts called a quantum (=Fixed quantity of energy).

• The relationship between energy and frequency is

where h is Planck’s constant (6.626 10-34 J.s).• There is an important relationship between energy and

wavelength of radiation:

Quantized Energy and Quantized Energy and PhotonsPhotons

hE

/hcE

Prentice Hall © 2003Chapter 6

The Photoelectric Effect and Photons• Einstein assumed that light traveled in energy packets

(stream of particles) called photons.• The energy of one photon:

Quantized Energy and Quantized Energy and PhotonsPhotons

hE

Photoelectric effect

• Must employ light of sufficient frequency (E=h)• Number of electrons ejected proportional to intensity

– More intense more photons

– Thus more electrons ejected

• Energy ejected electrons not proportional to intensity– h

Prentice Hall © 2003Chapter 6

• Rutherford assumed the electrons orbited the nucleus analogous to planets around the sun.

• However, a charged particle moving in a circular path should lose energy.

• This means that the atom should be unstable according to Rutherford’s theory.

• Bohr noted the line spectra of certain elements and assumed the electrons were confined to specific energy states. These were called orbits.

Line Spectra and the Bohr Line Spectra and the Bohr ModelModel


Bohr Model• Since the energy states are quantized, the light emitted

from excited atoms must be quantized and appear as line spectra.

• After lots of math, Bohr showed that

where n is the principal quantum number (i.e., n = 1, 2, 3, … and nothing else).


218 1

J 1018.2n

E

Bohr Model

Line Spectra and Line Spectra and the Bohr Modelthe Bohr Model


Bohr Model• The first orbit in the Bohr model has n = 1, is closest to

the nucleus, and has negative energy by convention.• The furthest orbit in the Bohr model has n close to

infinity and corresponds to zero energy.• Electrons in the Bohr model can only move between

orbits by absorbing and emitting energy in quanta (h).• The amount of energy absorbed or emitted on movement

between states is given by


hEEE if


Bohr Model• We can show that

• When ni > nf, energy is emitted.

• When nf > ni, energy is absorbed


2218 11

J 1018.2if nn

hchE


• Knowing that light has a particle nature, it seems reasonable to ask if matter has a wave nature.

• Using Einstein’s and Planck’s equations, de Broglie showed:

• The momentum, mv, is a particle property, whereas is a wave property.

• de Broglie summarized the concepts of waves and particles, with noticeable effects if the objects are small.

The Wave Behavior of The Wave Behavior of MatterMatter

mvh


The Uncertainty Principle• Heisenberg’s Uncertainty Principle: on the mass scale

of atomic particles, we cannot determine exactly the position, direction of motion, and speed simultaneously.

• For electrons: we cannot determine their momentum and position simultaneously.

• If x is the uncertainty in position and mv is the uncertainty in momentum, then

The Wave Behavior of The Wave Behavior of MatterMatter

4·

hmvx


• Schrödinger proposed an equation that contains both wave and particle terms.

• Solving the equation leads to wave functions. • The wave function gives the shape of the electronic

orbital.• The square of the wave function, ψ2 gives the probability

of finding the electron,• that is, gives the electron density for the atom.

Quantum Mechanics and Quantum Mechanics and Atomic OrbitalsAtomic Orbitals


Orbitals and Quantum Numbers• If we solve the Schrödinger equation, we get wave

functions and energies for the wave functions.• We call wave functions orbitals.• Keep in mind that orbital in the quantum-mechanical

model bears no resemblance to an orbit in the Bohr model. An orbit is an electron’s path around the nucleus whereas, an orbital is a mathematical function with no direct physical meaning.



An atomic orbital is specified by 3 quantum numbers:1. Principal Quantum Number, n, related to the size of the orbital. This is the

same as Bohr’s n. As n becomes larger, the atom becomes larger and the electron is further from the nucleus. The higher the value of n, the higher the energy level.

2. Orbital Angular Momentum (or, Azimuthal) Quantum Number, l. (related to shape). This quantum number depends on the value of n. The number of l value= the number of n value. The values of l begin at 0 and increase to (n - 1). We usually use letters for l (s, p, d and f for l = 0, 1, 2, and 3, respectively). Usually we refer to the s, p, d and f-orbitals.

3. Magnetic Quantum Number, ml related to orientation in space. This quantum number depends on l. The magnetic quantum number has integral values between -l and +l. The number of possible ml values equals the number of orbitals, which is 2l+1 for a given l value.




• Shell/level – The atom’s energy shells or levels, are given by the value n. The smaller the value of n, the lower the energy level and closer to the nucleus.

• Subshell/sublevel – The atom’s shells/levels contain subshells/sublevels, which designate the orbital shape. Each subshell has a letter designation:

• l = 0 is an s subshell, l = 1 is a p subshell l = 2 is a d subshell, l = 3 is an f subshell. • The letters derived from the names of spectroscopic lines: Sharp, principal, diffuse and fundamental.

Subshells are named by joining the n value and the letter designation. For example, n=2 and l=0 is called 2s subshell.



• Orbital – Each allowed combination of n, l and ml values specifies one of the atom’s orbitals to describe the shape, size and the spatial orientation.

• The value of n = the number of possible l values (an integer from 0 to n-1).

So, when n=2, l will have only two values, 0 and 1.

• The number of orbitals in each subshell is 2l+1 for a given l value. One s orbital (l=0), 3 p orbitals (l=1) and 5 d orbitals (l=2) and 7 f orbitals (l=3).

• Again, the total number of orbitals = n2.



• How many orbitals exist for n = 3?• For n = 3, l will have 3 values, i.e., 0, 1 and 2.

• For l = 0, ml will have 0 value (as -l, 0 and +l)

• For l = 1, ml will have 3 values (-1, 0 and +1)

• For l = 2, ml will have 5 values, -2 through 0 to +2. (i.e., -2, -1, 0 ,+1 and +2).

• There are 9 ml values which means 9 orbitals. In other words, n2=32=9.


Orbitals and Quantum Numbers



Orbitals and Quantum Numbers• Orbitals can be ranked in terms of energy to yield an

Aufbau diagram.• Note that the following Aufbau diagram is for a single

electron system.• As n increases, note that the spacing between energy

levels becomes smaller.



Orbitals and Quantum Numbers

Quantum Quantum Mechanics and Mechanics and

Atomic Atomic OrbitalsOrbitals


The s-Orbitals

Representations of Representations of OrbitalsOrbitals


The p, d and f -Orbitals

• There are three p-orbitals px, py, and pz.

• The three p-orbitals lie along the x-, y- and z- axes.

• The letters correspond to allowed values of ml of -1, 0, and +1.

• The orbitals are dumbbell shaped.• As n increases, the p-orbitals get larger.• There are five d and seven f-orbitals.



The p-Orbitals



The d-Orbitals



Many-Electron AtomsMany-Electron Atoms

Electron Spin and the Pauli Exclusion Principle

• Spin is an intrinsic property of an electron and this is another quantum number besides the other three. This is not a property of the orbitals.

• A beam of atoms was passed through a slit and into a magnetic field and the atoms were then detected.

• Two spots were found: one with the electrons spinning in one direction and one with the electrons spinning in the opposite direction.




• Since electron spin is quantized, we define ms = spin quantum number = ½.

• Pauli’s Exclusions Principle:: no two electrons can have the same set of 4 quantum numbers.

– Therefore, two electrons in the same orbital must have opposite spins.


Orbitals and Their Energies• The Aufbau diagram looks slightly different for many-

electron systems.

Many-Electron Atoms Many-Electron Atoms

Many-Electron Atoms Many-Electron Atoms

Orbitals and Their Energies


Hund’s Rule• Electron configurations tell us in which orbitals the

electrons for an element are located.• Hund’s rule: When more than one orbital has the same

energy, electrons occupy separate orbitals and do so with parallel spins.

• Electrons fill orbitals starting with lowest n and moving upwards;

• No two electrons can fill one orbital with the same spin (Pauli).

Electron ConfigurationsElectron Configurations


Condensed Electron Configurations• Neon completes the 2p subshell.• Sodium marks the beginning of a new row. So, we write

the condensed electron configuration for sodium as –

Na: [Ne] 3s1

[Ne] represents the electron configuration of neon.• Inner (core) electrons: electrons in [Noble Gas].They

fill all the lower energy levels of an atom.• Valence electrons: electrons outside of [Noble Gas].



Transition Metals• After Ar the d orbitals begin to fill.• After the 3d orbitals are full, the 4p orbitals being to fill.• Transition metals: elements in which the d electrons

are the valence electrons.



Lanthanides and Actinides• From Ce onwards the 4f orbitals begin to fill.• Note: La: [Xe]6s25d14f0

• Elements Ce - Lu have the 4f orbitals filled and are called lanthanides or rare earth elements.

• Elements Th - Lr have the 5f orbitals filled and are called actinides.

• Most actinides are not found in nature.



• The periodic table can be used as a guide for electron configurations.

• The period number is the value of n.• Groups 1A and 2A have the s-orbital filled.• Groups 3A - 8A have the p-orbital filled.• Groups 3B - 2B have the d-orbital filled.• The lanthanides and actinides have the f-orbital filled.

Electron Configurations Electron Configurations and the Periodic Tableand the Periodic Table


Molecular OrbitalsMolecular Orbitals

Molecular Orbital (MO) Theory.• Just as electrons in atoms are found in atomic orbitals,

electrons in molecules are found in molecular orbitals.



• Molecular orbitals:• each contain a maximum of two electrons;

• have definite energies;

• can be visualized with contour diagrams;

• are associated with an entire molecule.

The Hydrogen Molecule• When two AOs overlap, two MOs form.



The Hydrogen Molecule• Therefore, 1s (H) + 1s (H) must result in two MOs for

H2:

• one has electron density between nuclei (bonding MO);

• one has little electron density between nuclei (antibonding MO).

• MOs resulting from s orbitals are MOs. (bonding) MO is lower energy than * (antibonding)

MO.



The Hydrogen Molecule



The Hydrogen Molecule• Energy level diagram or MO diagram shows the

energies and electrons in an orbital.• The total number of electrons in all atoms are placed in

the MOs starting from lowest energy (1s) and ending when you run out of electrons.• Note that electrons in MOs have opposite spins.

• H2 has two bonding electrons.

• He2 has two bonding electrons and two antibonding electrons.



The Hydrogen Molecule



Bond Order• Define

• Bond order = 1 for single bond.

• Bond order = 2 for double bond.

• Bond order = 3 for triple bond.

• Fractional bond orders are possible.

• For H2

• Therefore, H2 has a single bond.

electrons gantibondin-electrons bondingorder Bond21

102order Bond21



Bond Order

• For He2

• Therefore He2 is not a stable molecule

022order Bond21


Second-Row Diatomic Second-Row Diatomic MoleculesMolecules

• We look at homonuclear diatomic molecules (e.g. Li2, Be2, B2 etc.).

• AOs combine according to the following rules:• The number of MOs = number of AOs;

• AOs of similar energy combine;

• As overlap increases, the energy of the MO decreases;

• Pauli: each MO has at most two electrons;• Hund: for degenerate orbitals, each MO is first occupied

singly.



Molecular Orbitals for Li2 and Be2

• Each 1s orbital combines with another 1s orbital to give one 1s and one *

1s orbital, both of which are occupied (since Li and Be have 1s2 electron configurations).

• Each 2s orbital combines with another 2s orbital, two give one 2s and one *

2s orbital.

• The energies of the 1s and 2s orbitals are sufficiently different so that there is no cross-mixing of orbitals (i.e. we do not get 1s + 2s).




• There are a total of 6 electrons in Li2:

• 2 electrons in 1s;

• 2 electrons in *1s;

• 2 electrons in 2s; and

• 0 electrons in *2s.

• Since the 1s AOs are completely filled, the 1s and *1s are filled. We generally ignore core electrons in MO diagrams.

124order Bond21




• There are a total of 8 electrons in Be2:

• 2 electrons in 1s;

• 2 electrons in *1s;

• 2 electrons in 2s; and

• 2 electrons in *2s.

• Since the bond order is zero, Be2 does not exist.

044order Bond21



Molecular Orbitals from 2p Atomic Orbitals

• There are two ways in which two p orbitals overlap:• end-on so that the resulting MO has electron density on the

axis between nuclei (i.e. type orbital);

• sideways so that the resulting MO has electron density above and below the axis between nuclei (i.e. type orbital).



Molecular Orbitals from 2p Atomic Orbitals

• The six p-orbitals (two sets of 3) must give rise to 6 MOs: , *, , *, , and *.

• Therefore there is a maximum of 2 bonds that can come from p-orbitals.

• The relative energies of these six orbitals can change.

Molecular Orbitals from

2p Atomic Orbitals



Configurations for B2 Through Ne2

• 2s Orbitals are lower in energy than 2p orbitals so 2s orbitals are lower in energy than 2p orbitals.

• There is greater overlap between 2pz orbitals (they point directly towards one another) so the 2p is MO is lower in energy than the 2p orbitals.

• There is greater overlap between 2pz orbitals so the *2p

is MO is higher in energy than the *2p orbitals.

• The 2p and *2p orbitals are doubly degenerate.




• As the atomic number decreases, it becomes more likely that a 2s orbital on one atom can interact with the 2p orbital on the other.• As the 2s-2p interaction increases, the 2s MO lowers in

energy and the 2p orbital increases in energy.

• For B2, C2 and N2 the 2p orbital is higher in energy than the 2p.

• For O2, F2 and Ne2 the 2p orbital is lower in energy than the 2p.




• Once the relative orbital energies are known, we add the required number of electrons to the MOs, taking into account Pauli’s exclusion principle and Hund’s rule.• As bond order increases, bond length decreases.

• As bond order increases, bond energy increases.



Electron Configurations and Molecular Properties

• The MO diagram for O2 predicts both paramagnetism and the double bond (bond order = 2).


• Chemical bond: attractive force holding two or more atoms together.

• Covalent bond results from sharing electrons between the atoms. Usually found between nonmetals.

• Ionic bond results from the transfer of electrons from a metal to a nonmetal.

• Metallic bond: attractive force holding pure metals together.

Chemical Bonds, and the Chemical Bonds, and the Octet RuleOctet Rule


The Octet Rule• All noble gases except He has an s2p6 configuration. • Octet rule: atoms tend to gain, lose, or share electrons

until they are surrounded by 8 valence electrons (4 electron pairs).

• Caution: there are many exceptions to the octet rule.

Chemical Bonds and the Chemical Bonds and the Octet RuleOctet Rule


Consider the reaction between sodium and chlorine:Na(s) + ½Cl2(g) NaCl(s) Hºf = -410.9 kJ

Ionic BondingIonic Bonding


• The reaction is violently exothermic.

• We infer that the NaCl is more stable than its constituent elements. Why?

• Na has lost an electron to become Na+ and chlorine has gained the electron to become Cl. Note: Na+ has an Ne electron configuration and Cl has an Ar configuration.

• That is, both Na+ and Cl have an octet of electrons surrounding the central ion.



• NaCl forms a very regular structure in which each Na+ ion is surrounded by 6 Cl ions.

• Similarly, each Cl ion is surrounded by six Na+ ions.• There is a regular arrangement of Na+ and Cl in 3D.• Note that the ions are packed as closely as possible.• Note that it is not easy to find a molecular formula to

describe the ionic lattice.




Energetics of Ionic Bond Formation• The formation of Na+(g) and Cl(g) from Na(g) and Cl(g)

is endothermic.• Why is the formation of Na(s) exothermic?• The reaction NaCl(s) Na+(g) + Cl(g) is endothermic

(H = +788 kJ/mol).• The formation of a crystal lattice from the ions in the gas

phase is exothermic:

Na+(g) + Cl(g) NaCl(s) H = 788 kJ/mol



Energetics of Ionic Bond Formation• Lattice energy: the energy required to completely

separate an ionic solid into its gaseous ions.• Lattice energy depends on the charges on the ions and the

sizes of the ions:

is a constant (8.99 x 10 9 J·m/C2), Q1 and Q2 are the charges on the ions, and d is the distance between ions.

dQQ

El21



Energetics of Ionic Bond Formation• Lattice energy increases as

• The charges on the ions increase• The distance between the ions decreases.


Ionic BondingIonic BondingElectron Configurations of Ions of the

Representative Elements• These are derived from the electron configuration of

elements with the required number of electrons added or removed from the most accessible orbital.

• Electron configurations can predict stable ion formation:• Mg: [Ne]3s2

• Mg+: [Ne]3s1 not stable

• Mg2+: [Ne] stable

• Cl: [Ne]3s23p5

• Cl: [Ne]3s23p6 = [Ar] stable


Covalent BondingCovalent Bonding

• When two similar atoms bond, none of them wants to lose or gain an electron to form an octet.

• When similar atoms bond, they share pairs of electrons to each obtain an octet.

• Each pair of shared electrons constitutes one chemical bond.

• Example: H + H H2 has electrons on a line connecting the two H nuclei.


Covalent BondingCovalent Bonding


Strengths of Covalent Strengths of Covalent BondsBonds

• The energy required to break a covalent bond is called the bond dissociation enthalpy, D. That is, for the Cl2 molecule, D(Cl-Cl) is given by H for the reaction:

Cl2(g) 2Cl(g).• When more than one bond is broken:

CH4(g) C(g) + 4H(g) H = 1660 kJ• the bond enthalpy is a fraction of H for the

atomization reaction:D(C-H) = ¼H = ¼(1660 kJ) = 415 kJ.

• Bond enthalpies can either be positive or negative.



Bond Enthalpies and the Enthalpies of Reactions

• We can use bond enthalpies to calculate the enthalpy for a chemical reaction.

• We recognize that in any chemical reaction bonds need to be broken and then new bonds get formed.

• The enthalpy of the reaction is given by the sum of bond enthalpies for bonds broken minus the sum of bond enthalpies for bonds formed.




• Mathematically, if Hrxn is the enthalpy for a reaction, then

• We illustrate the concept with the reaction between methane, CH4, and chlorine:

CH4(g) + Cl2(g) CH3Cl(g) + HCl(g) Hrxn = ?

formed bondsbroken bonds DDHrxn

Strengths of Covalent BondsStrengths of Covalent Bonds




• In this reaction one C-H bond and one Cl-Cl bond gets broken while one C-Cl bond and one H-Cl bond gets formed.

• The overall reaction is exothermic which means than the bonds formed are stronger than the bonds broken.

• The above result is consistent with Hess’s law.



Bond Enthalpy and Bond Length• We know that multiple bonds are shorter than single

bonds.• We can show that multiple bonds are stronger than

single bonds.• As the number of bonds between atoms increases, the

atoms are held closer and more tightly together.

Strengths of Covalent BondsStrengths of Covalent Bonds


Bond Polarity and Bond Polarity and ElectronegativityElectronegativity

• In a covalent bond, electrons are shared.• Sharing of electrons to form a covalent bond does not

imply equal sharing of those electrons.• There are some covalent bonds in which the electrons are

located closer to one atom than the other.• Unequal sharing of electrons results in polar bonds.



Electronegativity• Electronegativity: The ability of one atoms in a

molecule to attract electrons to itself.• Pauling set electronegativities on a scale from 0.7 (Cs) to

4.0 (F).• Electronegativity increases

• across a period and

• down a group.

Bond Polarity and ElectronegativityBond Polarity and Electronegativity

Electronegativity



Electronegativity and Bond Polarity• Difference in electronegativity is a gauge of bond

polarity:• electronegativity differences around 0 result in non-polar

covalent bonds (equal or almost equal sharing of electrons);• electronegativity differences around 2 result in polar covalent

bonds (unequal sharing of electrons);• electronegativity differences around 3 result in ionic bonds

(transfer of electrons).



Electronegativity and Bond Polarity• There is no sharp distinction between bonding types.• The positive end (or pole) in a polar bond is represented + and the negative pole -.



Dipole Moments• Consider HF:

• The difference in electronegativity leads to a polar bond.• There is more electron density on F than on H.• Since there are two different “ends” of the molecule, we call

HF a dipole.

• Dipole moment, , is the magnitude of the dipole:

where Q is the magnitude of the charges.• Dipole moments are measured in debyes, D.

Qr


Physical Properties of Metals• Important physical properties of pure metals: malleable,

ductile, good conductors, and feel cold.• Most metals are solids with the atoms in a close packed

arrangement.• In Cu, each atom is surrounded by 12 neighbors.• There are not enough electrons for the metal atoms to be

covalently bonded to each other.

Metallic BondingMetallic Bonding


Electron-Sea Model of Metallic Bonding• We use a delocalized model for electrons in a metal.

– The metal nuclei are seen to exist in a sea of electrons.

– No electrons are localized between any two metal atoms.

– Therefore, the electrons can flow freely through the metal.

– Without any definite bonds, the metals are easy to deform (and are malleable and ductile).

• Problems with the electron sea model:– As the number of electrons increase, the strength of bonding

should increase and the melting point should increase.



Electron-Sea Model of Metallic Bonding– But: group 6B metals have the highest melting points (center of

the transition metals).


Molecular-Orbital Model for Metals• Delocalized bonding requires the atomic orbitals on one

atom to interact with atomic orbitals on neighboring atoms.

• Example: graphite electrons are delocalized over a whole plane, benzene molecules have electrons delocalized over a ring.

• Recall: the number of molecular orbitals is equal to the number of atomic orbitals.

• In metals there is a very large number of orbitals.



Molecular-Orbital Model for Metals• As the number of orbitals increase, their energy spacing

decreases and they band together.• The number of electrons do not completely fill the band

of orbitals.• Therefore, electrons can be promoted to unoccupied

energy bands.• Since the energy differences between orbitals are small,

the promotion of electrons occurs at low energy costs.


Molecular-Orbital Model for Metals



Molecular-Orbital Model for Metals• As we move across the transition metal series, the

antibonding band starts becoming filled.• Therefore, the first half of the transition metal series have

only bonding-bonding interactions, the second half has bonding-antibonding interactions.

• We expect the middle of the transition metal series to have the highest melting points.

• The energy gap between bands is called the band gap.


Ketelaar TriangleKetelaar Triangle


Covalent

Ionic

Metallic


Exceptions to the Octet Exceptions to the Octet RuleRule

• There are three classes of exceptions to the octet rule:• Molecules with an odd number of electrons;

• Molecules in which one atom has less than an octet;

• Molecules in which one atom has more than an octet.

Odd Number of Electrons

• Few examples. Generally molecules such as ClO2, NO, and NO2 have an odd number of electrons.

N O N O



Less than an Octet• Relatively rare.• Molecules with less than an octet are typical for

compounds of Groups 1A, 2A, and 3A.

• Most typical example is BF3.

• Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.



More than an Octet• This is the largest class of exceptions.• Atoms from the 3rd period onwards can accommodate

more than an octet.• Beyond the third period, the d-orbitals are low enough in

energy to participate in bonding and accept the extra electron density.


• Physical properties of substances understood in terms of kinetic molecular theory:– Gases are highly compressible, assumes shape and volume of

container: • Gas molecules are far apart and do not interact much with each

other.

– Liquids are almost incompressible, assume the shape but not the volume of container:

• Liquids molecules are held closer together than gas molecules, but not so rigidly that the molecules cannot slide past each other.

A Molecular Comparison A Molecular Comparison of Liquids and Solidsof Liquids and Solids


– Solids are incompressible and have a definite shape and volume:

• Solid molecules are packed closely together. The molecules are so rigidly packed that they cannot easily slide past each other.



• Converting a gas into a liquid or solid requires the molecules to get closer to each other:– cool or compress.

• Converting a solid into a liquid or gas requires the molecules to move further apart: – heat or reduce pressure.

• The forces holding solids and liquids together are called intermolecular forces.



• The covalent bond holding a molecule together is an intramolecular forces.

• The attraction between molecules is an intermolecular force.

• Intermolecular forces are much weaker than intramolecular forces (e.g. 16 kJ/mol vs. 431 kJ/mol for HCl).

• When a substance melts or boils the intermolecular forces are broken (not the covalent bonds).

Intermolecular ForcesIntermolecular Forces


Ion-Dipole Forces• Interaction between an ion and a dipole (e.g. water).• Strongest of all intermolecular forces.



Dipole-Dipole Forces• Dipole-dipole forces exist between neutral polar

molecules.• Polar molecules need to be close together.• Weaker than ion-dipole forces.• There is a mix of attractive and repulsive dipole-dipole

forces as the molecules tumble.• If two molecules have about the same mass and size, then

dipole-dipole forces increase with increasing polarity.


Dipole-Dipole Forces



London Dispersion Forces• Weakest of all intermolecular forces.• It is possible for two adjacent neutral molecules to affect

each other.• The nucleus of one molecule (or atom) attracts the

electrons of the adjacent molecule (or atom).• For an instant, the electron clouds become distorted.• In that instant a dipole is formed (called an instantaneous

dipole).



London Dispersion Forces• One instantaneous dipole can induce another

instantaneous dipole in an adjacent molecule (or atom).• The forces between instantaneous dipoles are called

London dispersion forces.



London Dispersion Forces• Polarizability is the ease with which an electron cloud

can be deformed.• The larger the molecule (the greater the number of

electrons) the more polarizable.• London dispersion forces increase as molecular weight

increases.• London dispersion forces exist between all molecules.• London dispersion forces depend on the shape of the

molecule.



London Dispersion Forces• The greater the surface area available for contact, the

greater the dispersion forces.• London dispersion forces between spherical molecules

are lower than between sausage-like molecules.


London Dispersion Forces



Hydrogen Bonding• Special case of dipole-dipole forces.• By experiments: boiling points of compounds with H-F,

H-O, and H-N bonds are abnormally high.• Intermolecular forces are abnormally strong.



Hydrogen Bonding• H-bonding requires H bonded to an electronegative

element (most important for compounds of F, O, and N).– Electrons in the H-X (X = electronegative element) lie much

closer to X than H.

– H has only one electron, so in the H-X bond, the + H presents an almost bare proton to the - X.

– Therefore, H-bonds are strong.


Hydrogen Bonding


Hydrogen Bonding• Hydrogen bonds are responsible for:

– Ice Floating• Solids are usually more closely packed than liquids;

• Therefore, solids are more dense than liquids.

• Ice is ordered with an open structure to optimize H-bonding.

• Therefore, ice is less dense than water.

• In water the H-O bond length is 1.0 Å.

• The O…H hydrogen bond length is 1.8 Å.

• Ice has waters arranged in an open, regular hexagon.

• Each + H points towards a lone pair on O.


Hydrogen Bonding


Chapter 1

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